Thermochemistry

Thermochemistry

Energy • Energy is necessary for all life. • The study of energy and it transformations is known as thermodynamics. • In chemical reactions we study that aspect of thermodynamics that involves the relationships between chemical reactions and energy changes involving heat. This relationship is called thermochemistry.

The nature of energy • Energy is defined as the capacity to do work or to transfer heat. • Work is the force required to move an object against. It is the energy required to overcome friction and inertia. • Heat is the energy used to cause the temperature of an object to increase.

Kinetic Energy • Kinetic energy is the energy of motion • Ek = ½ mv2 • The kinetic energy of an object increases as the mass increases and as the square of the velocity increases. • A 1000 kg (2200 pounds) car traveling at 10 m/s (22.5 mph) has a kinetic energy of 50kJ. The same car traveling at 20 m/s (45 mph) has a kinetic energy of 200 KJ.

Potential Energy • Potential energy is the energy of position, it arises when a force operates on an object. • The two potential energies we are familiar with are gravitational and electrostatic. • Ep = mgh • Where m is the mass of the object, h is the height of the object relative to a reference height and g is the gravitational constant of 9.8 m/s2

Potential Energy (cont) • As chemists we don’t concern ourselves with the potential energy of gravitation due to the small size of the objects concerned. • Very concerned with electrostatic potential. • The electrostatic potential arises from the interactions between charged particles. • Eel = kQ1Q2/d • Where k is a constant of proportionality = 8.99 x 109 J-m/C2

Potential Energy (cont) • The chemical energy of molecules is due to the potential energy stored in the arrangements of their atoms.

Units of Energy • The SI unit of energy is the joule, J • The car from the previous example – • A 1000 kg (2200 pounds) car traveling at 10 m/s (22.5 mph) • Ek = ½ mv2 • = ½ (1000 kg)(10 m/s)2 • = ½ (1000 kg)(100 m2/s2) • = 50,000 kg-m2/s2 • = 50 kJ

Units of Energy • The calorie was originally defined as the amount of energy required to raise the temperature of 1 g of water from 14.5 oC to 15.5 oC • 1 cal = 4.184 J (exactly) • Side note the calorie in food units is equal to 1 kcal.

System and Surroundings • We can not account for the entire universe when performing our analysis. • We focus on a small portion of the universe and that portion is called the system, everything else is the surroundings. • When we study the energy change that accompanies a chemical reaction, the reactants and products constitute the system.

Systems • A closed system can exchange energy but not matter with the surroundings. • A car’s cylinder during ignition is a closed system. • An open system can exchange energy and matter with the surroundings. • A open beaker with boiling water • An isolated system can exchange neither energy nor matter with the surroundings. • An insulated thermos approximates an isolated system.

Transferring Energy: Work & Heat • Energy is transferred between systems and surroundings via heat or work or both. • Energy used to cause an object to move is called work. • W = F x d • We perform work when we lift an object against the force of gravity. • Heat is the energy transferred from a hotter object to a colder one. Heat always flows from hotter to colder objects.

The First Law of Thermodynamcis • Energy is conserved – energy can neither be created or destroyed only converted from one form to another. • The internal energy of a system is the sum of all the kinetic energy and potential energies of all its components.

Internal Energy • The internal energy includes translational energies, vibrational energies, rotational energies of every atom in the system. • The numerical value of the internal energy of a system is generally not known to a high degree of accuracy. (Why?) • We can calculate the change in internal energy. • ΔE = Efinal – Einitial

Internal Energy • ΔE contains three parts • A number • A unit • A direction • A negative value of ΔE indicates the system has lost energy to its surroundings • A positive value of ΔE indicates the system has gained energy from its surroundings • It is important to note that energy changes are from the point of view of the system not the surroundings.

Internal Energy (cont) • In a chemical reaction, the initial state of the system refers to the reactants. • The final state of the system refers to the products. • When hydrogen and oxygen form water at a given temperature, the system loses energy to the surroundings, thus the internal energy of the products is less than that of the reactants.

ΔE – heat and work • The internal energy of a system changes in magnitude as heat is added to or removed from the system or as work is done on it or by it. • ΔE = q + w • When heat is added to a system or work is done on a system, its internal energy increases.

Endothermic and Exothermic Processes • When a process occurs in which the system absorbs heat, the process is called endothermic. • During an endothermic process heat flows into the system from its surroundings • When a process occurs in which the system emits heat, the process is called exothermic. • During an exothermic process heat flows from the system to its surroundings.

State functions • The value of a state function depends only on the present state of the system, not on the path the system took to reach that state. • Because E is a state function, ΔE, depends only on the initial and final states of the system, not on how the change occurs. • E is a state function however q and w are not.

Enthalpy • Even though most of the chemical reactions we will examine occur at a constant atmospheric pressure, work is still being performed. (w = F x d) • As an example dissolving Zn in and acid release H2 gas. The expansion of the gas against atmospheric pressure is work. • The work involved in the expansion or compression of a gas is called pressure-volume work. • w = -P ΔV

Enthalpy (cont) • Enthalpy is a thermodynamic function that accounts for heat flow in a process occurring at constant pressure when the only work performed is P-V work. • Enthalpy is denoted by the symbol H • H = E + PV • The change in enthalpy of a system equals the heat gained or lost at constant pressure. • p 176

Enthalpies of reactions • Since ΔH = Hproducts – Hreactants the enthalpy change that accompanies a reaction is called the enthalpy of reaction or heat of reaction. • When exactly 2 moles of H2 reacts with exactly 1 mole of O2 to form 2 moles of H2O at a constant pressure, the system releases 483.6 kJ of heat. • 2H2 + O2 2H2O(g) ΔH = -483.6 kJ • Examine energy state diagram p 177

Thermochemical equations and enthalpy diagrams • Enthalpy is an extensive property. • The magnitude of ΔH is directly proportional to the amount of reactant consumed in the process. • CH4(g) + O2(g)  CO2(g) + 2H2O(l) ΔH = -890 kJ • If two moles of methane are consumed then • ΔH = -1780 kJ • The enthalpy change for a reaction is equal in magnitude, but opposite in sign, to the ΔH of the reverse reaction. • CO2 (g) + 2H2O(l)  CH4(g) + O2(g) ΔH = 890 kJ

Thermochemical equations and enthalpy diagrams (cont) • The enthalpy change for a reaction depends on the state of the reactants. • If the products in the combustion of methane were gaseous H2O instead of liquid H2O ΔH would be -802 kJ instead of – 890 kJ. • It is necessary to specify the state of the reactants and products.

Calorimetry • The value of ΔH can be determined experimentally via a calorimeter. • The measurement of heat flow is calorimetry. • We determine the magnitude of the heat flow by measuring the magnitude of the temperature change.

Statements of the Obvious • The more heat an object gains, the hotter it gets. • All substances change temperature when they are heated. • But less obvious is that the magnitude of the change with a given heat varies from substance to substance.

Heat Capacity and Specific Heat • The temperature change experienced by an object when it absorbs certain amount of heat is determined by its heat capacity, C. • The heat capacity of an object is the amount of heat required to raise its temperature by 1oC. • The heat capacity of a mole of a substance is it molar heat capacity. The heat capacity of one gram of a substance is its specific heat.

Specific Heat • Specific heat = (quantity of heat transferred) • (gram of substance)(temp. change) • Cs = q/(m x ΔT) • Exercise 5.5 p 181

Calorimetry • Constant-Pressure calorimetry is the process of measuring the heat transfer at constant, usually atmospheric, pressure. • Constant-Volume calorimetry is the process of measuring heat flow at constant volume in a device called a bomb calorimetery. (p 183)

Hess’s Law • Many enthalpies of reaction have been tabulated and from those tabulations it is possible to calculate the enthalpy of a reaction without the need for a calorimeter. • Hess’s law states that if a reaction is carried out in a series of steps, ΔH for the overall reaction will equal the sum of the enthalpies of reaction of the individual steps.

Enthalpies of Formation • Enthalpy of formation is the enthalpy or heat change during formation of a compound from its constituent elements. • The magnitude of any enthalpy change depends on the conditions of temperature, pressure and state. • Therefore to compare enthalpies we must define a set of conditions called the standard state. • STP – standard temperature and pressure

Standard enthalpy change • The standard enthalpy change of a reaction is defined as the enthalpy change when all reactants and products are in their standard states. • By definition, the standard enthalpy of formation of the most stable form of any element is zero.

Examples • Example 5.10 pg 190 • Using enthalpies of formation to calculate enthalpies of reaction

Foods • The vast majority of the energy our bodies consume comes from carbohydrates and fats. • C6H12O6 + 6O2 6CO2 + 6H2O ΔH = -2803 kJ • Tristearin is a typical fat, • 2C57H110O6(s) + 163O2(g)  114 CO2 + 110 H2O(l) ΔH = -75,520 kJ

Fuels • Fossil Fuels • Natural gas • Petroleum • Coal

Thermochemistry